2.6 AS 91166 Demonstrate understanding of chemical reactivity (4 credits)
NCEA PAST EXAM QUESTIONS on Rates of Reaction
The table below provides an overview of the specify factor affecting the Rate of a Reaction and the year that this question appeared in the NCEA Exam paper
(a) Hydrochloric acid was reacted with calcium carbonate in the form of marble chips (lumps) and powder (crushed marble chips) in an experiment to investigate factors affecting the rate of a chemical reaction.
(i) Identify the factor being investigated.
(ii) Explain why the hydrochloric acid would react faster with the powder.
(b) A clock reaction involves mixing solution X and solution Y with starch present. When the reaction is complete the solution turns blue-black in colour.
A student carried out this reaction between solution X and solution Y in a conical flask. Over time, the cross on the piece of paper under the flask disappeared when viewed from above.
The following experiments were carried out, and the times taken for the cross to disappear recorded.
Elaborate on why the reactions in Experiment 2 and Experiment 3 occur faster than the reaction in Experiment 1. In your answer, include the following words or terms: collisions activation energy temperature effective catalyst
When dilute hydrochloric acid, HCl(aq), is added to sodium thiosulfate, Na2S2O3(aq), in a conical flask, the following reaction occurs:
A pale yellow solid of sulfur, S(s), forms during the reaction. Over time, a cross on a piece of paper under the conical flask gradually disappears when viewed from above.
(a) List TWO ways that the rate of this reaction could be decreased.
(b) The following experiments were carried out, and the times taken for the cross to disappear recorded. The HCl(aq) was in excess in all of the experiments.
Analyse how the results of Experiment 2 and Experiment 3 compare to Experiment 1.
In your answer you should:
• identify the factor being changed and the effect it has on the reaction rate
• explain how the rate of reaction was affected, with reference to the collision of particles, and activation energy where appropriate.
Experiment 2 compared to Experiment 1:
Experiment 3 compared to Experiment 1:
The reaction between hydrochloric acid solution and sodium thiosulfate solution produces a precipitate of sulfur that makes the resulting solution go cloudy.
The reaction is carried out in a conical flask, which has been placed on a piece of paper with a cross, as shown in the diagram below.
The rate of this reaction can be measured by timing how long it takes for the solution to go so cloudy that the cross can no longer be seen through the solution.
(a) Describe how the rate of this reaction would change if the reaction was repeated:
(i) with a more diluted solution of sodium thiosulfate.
(ii) at a lower temperature.
(b) Explain how decreasing the concentration of sodium thiosulfate and decreasing the temperature of the reactants affects the reaction rate. In your answer, you should refer to:
particle collision
activation energy
the reaction between hydrochloric acid and sodium thiosulfate
(a) Zinc foil reacts with dilute sulfuric acid to produce zinc sulfate and hydrogen gas.
Zn(s) + H2SO4(aq) → ZnSO4(aq) + H2(g)
Two experiments were carried out using the same mass of zinc foil and excess sulfuric acid of concentrations 0.500 mol L–1 and 1.00 mol L–1. The reactions were carried out at 25°C. A graph showing the results of the two experiments is shown below.
(i) Write down the letter of the experiment which was carried out using the 1.00 mol L–1 sulfuric acid. Justify your answer in terms of the collision of particles.
(ii) Experiment A was repeated, using the same mass of zinc granules instead of zinc foil. The rate of reaction increased. Explain this effect on the reaction rate, in terms of the factor investigated and the collision of particles.
(iii) Experiment A was repeated again; this time the reaction was carried out at a temperature of 15°C. Discuss the effect of this temperature change.
Your answer must include reference to:
reaction rate
collision of particles
Activation Energy
(a) An experiment to investigate the rate of reaction between a 5 cm strip of magnesium ribbon and 20 mL of 1.00 mol L–1 hydrochloric acid was carried out at 25°C. The hydrochloric acid was in excess.
The experiment was repeated under different conditions as given below.
Complete the table below for EACH change, by:
stating the factor being investigated
describing how the reaction rate would be affected by using the word "increase" or "decrease" in your answer
(a) An experiment was carried out by reacting calcium carbonate with dilute hydrochloric acid.
The equation for this reaction can be represented by:
CaCO3(s) + 2HCl(aq) → CaCl2(aq) + CO2(g) + H2O(ℓ)
State THREE ways to increase the rate of this reaction.
(b) Three experiments were carried out by reacting magnesium metal with hydrochloric acid.
The equation for this reaction can be represented by:
Mg(s) + 2HCl(aq) → MgCl2(aq) + H2(g)
The beakers drawn below show particle representations of magnesium metal and hydrochloric acid just before the reaction is about to begin.
The rate of reaction can be followed by recording the volume of hydrogen, H2, produced. The graph below shows the volume of hydrogen produced during the initial part of the reaction.
Which beaker(s) corresponds to: LINE A of the graph and LINE B of the graph _
Explain your answers in terms of
particle collision
concentration of the particles present
(c) An experiment was carried out by reacting bromine, Br2, and methanoic acid, HCOOH, in a beaker. The equation for this reaction can be represented by:
The rate of reaction can be followed by measuring the time taken for the solution to go colourless.
A series of reactions were carried out and the time taken for the solution to go colourless was recorded. The results are shown below.
(i) Describe the trend shown by these results.
(ii) Discuss the reasons for this trend.
Your answer must include reference to
particle collision
reaction rate
Activation Energy
Hydrogen peroxide decomposes to form water and oxygen gas.2H2O2(aq) → 2H2O(l) + O2(g)
The rate of decomposition is changed with the addition of manganese dioxide, MnO2.
The rate of the reaction can be followed by recording the volume of oxygen produced. The graph below shows the volume of oxygen produced with, and without, manganese dioxide.
(a) Describe what happens to the rate of decomposition of hydrogen peroxide when manganese dioxide is added.
(b) A student suggested that manganese dioxide is a catalyst for the decomposition reaction.
Explain the role of a catalyst in changing the rate of a reaction.
Your answer must include reference to:
particle collision
Activation Energy
(c) One line on the graph above is divided into two time periods, A and B.
Discuss the change in reaction rate in terms of particle collisions during these two time periods, A and B.
An experiment was carried out by reacting zinc metal with excess dilute sulfuric acid. All the zinc was used up in the reaction. The equation for this reaction is:Zn(s) + H2SO4(aq) → ZnSO4(aq)+ H2(g)
(a) The experiment was repeated, but this time 2 mL water was also added to the beaker. Again the zinc was all used up. Explain why: when water is added, the rate of reaction will decrease.
(b) Small pieces of copper can act as a catalyst for the reaction between zinc and dilute sulfuric acid.
Discuss how this claim could be tested by experiment.
In an experiment, a sample of large marble chips (CaCO3) is added to 200 mL of dilute hydrochloric acid in an open conical flask. The reaction that occurs is shown below.
2HCl(aq) + CaCO3(s) → CaCl2(aq) + H2O(ℓ) + CO2(g)
As the carbon dioxide gas escapes from the flask, the total mass of the flask and contents decreases. The loss in mass is recorded at 5 minute intervals until the reaction has stopped. The experiment is repeated, using the same mass, but different sized, marble chips.
The results are shown in the table below.
(a) What factor affecting reaction rate is being investigated in this experiment?
(b) What conclusion can be made about the reaction rate using the data in the table?
(c) Explain the change in reaction rate that occurs, with reference to the collisions of particles.
(1) Hydrogen peroxide decomposes at room temperature (25'C) according to the following equation.2H2O2(aq) → 2H2O(l) + O2(g)
As the decomposition reaction occurs, bubbles of gas are produced.
(a) On addition of a very small amount of solid manganese dioxide, the rate at which the bubbles of gas are produced is increased so that rapid fizzing is observed. Further observation indicates that manganese dioxide remains after reaction has stopped. With reference to the collisions of particles, explain why the reaction rate has increased.
(b) Hydrogen peroxide is stored at a low temperature. Discuss this statement in terms of reaction rate.
(2) The concentration and pH of three acids, HA, HB and HC, are shown in the table below.
(a) A small piece of magnesium is added to a 20 mL sample of each of the acid. Name the acid that would be expected to react most rapidly with the magnesium. Explain why this acid will react the fastest.
The reaction between 20.0 mL of 0.500 mol L–1 hydrochloric acid and 20.0 mL of 0.250 mol L–1 sodium thiosulfate solution at room temperature (25°C) produces a precipitate of sulfur that makes the solution go cloudy after about 5 minutes.
(a) How would the time taken for the solution to go cloudy be affected if the reaction were carried out in a water bath at a temperature of 50°C?
(b) With reference to the collisions of particles, explain why the reaction is affected in this way.
Attempt all of the previous NCEA Rate of Reaction exam questions - shown above and available to download on a word document below. Ideally get your answer corrected by either your teacher or email them to the email link on the sidebar, we'll email you back with feedback in a couple of days. It is essential that you have personal feedback on your answers so as to ensure you have phrased your answer correctly and not left out any key points.
There is no set of Crystal Ball questions for Rates of Reaction, once you have completed all above you'll be very comfortable answering a question on rates of reaction.
NCEA EXAM QUESTION:Write the Equilibrium constant expression for the following reactions
NCEA EXAM QUESTION: Interpreting information using equilibrium constant values
1) For the reaction below, the values of Kc at different temperatures are shown in the table. N2(g) + 3H2(g) ⇄ 2NH3(g)
Use this information to determine whether the formation of NH3(g) is endothermic or exothermic. Justify your reasoning using equilibrium principles.
2) For the reaction below, the Kc value is 46.8 at 491°C H2(g) + I2(g) ⇄2HI(g) Calculate the concentration of HI(g), at equilibrium, at 491°C, if the concentration of H2(g) is 0.0190 mol L–1 and the concentration of I2(g) is 0.210 mol L–1.
The table below shows the value of the equilibrium constant, Kc at two different temperatures.
(i) Circle the species that will be in the highest concentration at 200°C. PCl5(g) or PCl3(g)
(ii) Explain your answer.
(iii) Calculate the concentration of PCl5 at equilibrium at 350°C, if the concentrations of PCl3 and Cl2 are both 0.352 mol L–1.
The reaction between sulfur dioxide gas and oxygen gas at a particular temperature to make sulfur trioxide gas can be represented as:2SO2(g) + O2(g) --> 2SO3(g)
For this equilibrium at this temperature, the value of Kc is 280.
The following three gases are mixed in a 1L container.
Justify whether the reaction would proceed in the forward OR reverse direction when the gases are mixed.
The following reaction can be used to produce gaseous methanol, CH3OH, from carbon monoxide and hydrogen.
CO(g) + 2H2(g) → CH3OH(g) ΔrH = –90.7 kJ mol–1
At 25°C, the equilibrium constant, Kc = 2.20 × 10–4. Explain what this indicates about the relative amounts of reactants and product at equilibrium.
Hydrogen gas and iodine vapour are placed in a sealed container at 445°C. These gases combine to form hydrogen iodide gas. This equilibrium can be represented by:
(a) Describe how an observer would know that the system had reached equilibrium.
(b) At 445°C, the equilibrium constant Kc = 49.5.
(i) Name the species that will be in the highest concentration at this temperature. You may name one, or more than one species in your answer.
(ii) Explain your answer.
(c) When the temperature of the equilibrium system is raised from 445°C to 1000°C (at constant pressure), the value of Kc decreases.
Use this information to determine whether the reaction between hydrogen and iodine is exothermic or endothermic. Justify your reasoning.
(a) Nitrogen monoxide gas reacts with oxygen gas to form nitrogen dioxide gas. The equilibrium reaction can be represented by:
2NO(g) + O2(g) → 2NO2(g)colourless colourless brown
At 230°C the equilibrium constant for this reaction has a value of 6.44 × 105.
(ii) State which gas will be in the highest concentration at 230°C. Explain your answer in terms of Kc and the colour seen.
Ag+(aq) + 2NH3(aq) → Ag(NH3)2+(aq)(i) At 25°C the value of Kc is 1.70 ´ 107. Name the species that would be present in the higher concentration in the equilibrium mixture at this temperature. Choose either Ag+(aq) or Ag(NH3)2+(aq) Justify your choice.2NO2(g) → 2NO(g) + O2(g)(ii) At 200°C the value of Kc is 1.10 ´ 10–5. Name the species that would be present in the higher concentration in the equilibrium mixture at this temperature. Choose either NO2(g) or NO(g). Justify your choice.
NCEA EXAM QUESTION: Describing and Explaining equilibrium constant reactionsNC
NCEA EXAM QUESTION: Describing and Explaining equilibrium constant reactions
The two reactions shown in the following table are both at equilibrium.
Compare and contrast the effect of increasing the pressure on both reactions, with reference to the equilibrium positions.
1) For each of the following changes applied to this system:
(i) State if the amount of chlorine gas, Cl2(g),would increase or decrease.
(ii) Justify your answers using equilibrium principles.
PCl3(g) is removed
The pressure is decreased.
2) When the temperature of the equilibrium system is increased from 200°C to 350°C (at constant pressure), the value of Kc increases, as shown in the table
Use this information to determine whether the decomposition of PCl5 is endothermic or exothermic. Justify your reasoning using equilibrium principles.
(a) Two oxides of nitrogen exist in an equilibrium system represented by the following equation. N2O4 is a colourless gas and NO2 is a brown gas.
N2O4(g) → 2NO2(g) ∆rH = + 57 kJ mol–1
For each of the following changes applied to this equilibrium system in (i) and (ii) below:
describe the expected observation
use equilibrium principles to explain this observation.
(i) The reaction mixture is cooled.
(i) The pressure is decreased.
The following reaction can be used to produce gaseous methanol, CH3OH, from carbon monoxide and hydrogen.
CO(g) + 2H2(g) → CH3OH(g) ΔrH = –90.7 kJ mol–1
Below are three possible factors that could be changed to alter the amount of methanol produced in this reaction.
Discuss the changes which could be made to increase the amount of methanol produced. Your answer should relate each of the factors in the table above to equilibrium principles.
The following equilibrium system is formed when hydrogen gas is mixed with iodine gas.
The reaction has a negative value for ΔrH. For each of the following changes applied to this system:
describe the expected observation
use equilibrium principles to discuss the reason for this observation.
(i)HI(g) is added.
(ii)The reaction mixture is cooled.
(iii) The pressure is increased.
(1) One step in the production of sulfuric acid involves forming sulfur trioxide from sulfur dioxide. The equilibrium reaction can be represented by
2SO2(g) + O2(g) → 2SO3(g) ΔrH = –196 kJ mol–1
(a) (i) Explain why a low temperature favours the formation of SO3(g).
(ii) The temperature that is actually used is approximately 450°C. However, this is not considered to be a
low temperature. Discuss why this temperature is used.
(b) (i) Describe another way of increasing the amount of SO3(g) present at equilibrium without adding any more reactants.
(ii) Explain why this will increase the amount of SO3(g) present at equilibrium.
(2) Carbon dioxide is added to drinks to make them fizzy. The following equilibria are involved:
The drink is fizzy when there is dissolved carbon dioxide, CO2(aq). The drink stops being fizzy when the carbon dioxide escapes from the drink as a gas.
Using equilibrium principles, discuss the changes that occur as a bottle containing fizzy drink is opened.
Your answer must include reference to:
equilibrium shift in Equation One and Equation Two
changes in the fizziness of the drink
any change in pH.
The following equilibrium system is formed when potassium thiocyanate solution is added to a solution of iron(III) nitrate. The reaction has a positive value for ΔrH
For each of the following changes applied to this system:
describe the expected observation
use equilibrium principles to discuss the reason for this observation.
(i) The reaction mixture is cooled.
(ii) Solid sodium fluoride is added to the reaction mixture. The fluoride ions react with Fe3+ ions.
(iii) Solid iron(III) chloride is added to the reaction mixture.
An equilibrium system is shown below.
3H2(g) + N2(g) --> 2NH3(g)
The pressure of the system is increased, while maintaining a constant temperature. The percentage of NH3 in the reaction mixture is recorded and graphed.
(b) On the above graph, identify the line that shows the correct relationship between the percentage of NH3 in the reaction mixture, and increasing pressure. Explain your answer by applying knowledge of equilibrium principles.
(2) An equilibrium system involving different species of cobalt(II) is shown in the equation below.
At room temperature (25°C) the equilibrium mixture is pink.
(a) Describe the expected observation when solid sodium chloride (NaCl) is added to the equilibrium mixture. Explain your answer.
(b) The enthalpy change (∆rH) for this reaction as written above, has a negative value.
Name the ion that would be present in the higher concentration when the equilibrium mixture is heated, choose either [CoCl4]2–(aq) or [Co(H2O)6]2+(aq) Explain your answer.
The following reaction is exothermic:
Both N2O5 and O2 are colourless gases and NO2 is a brown gas. A mixture of these gases exists at equilibrium and is observed as a brown colour.
(a) For each of the following changes applied to the equilibrium system, describe the expected observation and explain why this occurs.
(i) The mixture of gases is heated (at constant pressure).
(ii) The pressure is increased, by decreasing the volume of the container.
(a) The following equilibrium system is established when thiocyanate ions (SCN) are added to iron (III) ions (Fe3+). The resulting aqueous solution is a dark red colour. The equation representing the equilibrium system and the colours of each species involved are given below.
When iron (III) ions (Fe3+) are removed from the equilibrium mixture (by adding sodium fluoride), a colour change is observed. Describe the colour change you would expect to see and explain why it occurs.(ii)
The pressure of the system at equilibrium is increased (by decreasing the total volume of the system). Describe the effect of this change on the amount of NH3 in the system. Explain your answer.The percentage of NH3 present in equilibrium mixtures at different temperatures and at constant pressure is shown in the table below.
Justify whether the reaction in which NH3 is formed, is endothermic or exothermic.
There is no Crystal Ball Activity for this task as once you have completed all the above calculations you will be sweet with them. NCEA PAST EXAM QUESTIONS on writing equations on conjugate acid/base pair
(i) Complete the table below to show the conjugate acid-base pairs.
(ii) HPO42–(aq) is a species that can act as an acid or a base. Write equations for the reactions of HPO42– with water: one where it acts as an acid, and one where it acts as a base.
Complete the equations below to show how each species will react with water.
i) NH3 + H2O
ii) C2H5COO– + H2O
i) Identify one conjugate acid / base pair for the following reaction:
NH4+ + PO43– --> NH3 + HPO42–
ii) HSO4–(aq) is a species that can act as an acid or a base.
Write two equations for reactions of HSO4– with water: one equation where it acts as an acid, and one where it acts as a base.
Complete the equations below to show how each species will react with water to form an acidic solution.
i) NH4+(aq) + H2O -->
ii) HF(aq) + H2O -->
The bicarbonate ion, HCO3–, can both accept and donate hydrogen ions (protons). Write the equations below. When sodium bicarbonate, NaHCO3, dissolves in water the solution is basic.
Circle Reaction A or Reaction B to show which reaction predominates. Justify your answer.
i) Complete the table below to show the conjugate acid-base pairs.
ii) Circle the ion below that can act as both an acid and a base. Justify your choice CH3COO– HCO3–
iii) Hypochlorous acid is a weak acid. Complete the equation below to show the reaction of hypochlorous acid with water.
iii) A solution of sodium hypochlorite, NaOCl, is basic. Discuss the above statement, including appropriate chemical equation(s) in your answer.
Chickens make egg shell, CaCO3, using carbon dioxide gas from the air. The carbon dioxide forms carbonic acid (H2CO3), which then reacts to form the carbonate ions (CO32–) needed to make egg shell. Two equations showing part of this process are given below.
NCEA PAST EXAM QUESTIONS on Explaining properties of aqueous solutions
Some properties of three aqueous solutions A, B and C, of equal concentration are shown in the table below.
The labels of the three solutions have been removed. It is known that the solutions are NH3(aq), HCl(aq) and NH4Cl(aq). Use the information in the table above to identify each of the three solutions. Justify the identification of all three solutions. In your answer you should:
• refer to both pH and electrical conductivity of the solutions
• link your answers to appropriate chemical equations.
(a) Place the following solutions in order of increasing pH,
0.01 mol L–1 CH3COOH
0.01 mol L–1 HCl
0.1 mol L–1 HCl
0.1 mol L–1 NaOH
Justify your order above in terms of: proton transfer and relative concentrations of [H3O+] and [OH–], linked to the pH of the solution. You should include equations in your answer.
Three aqueous solutions, of equal concentration, have the following pH values
Compare and contrast both the strength and electrical conductivity of these aqueous solutions. Include appropriate equations in your answer.
i) The concentration and pH of two acids HA and HB are shown in the table below. Identify which one of these acids is weaker and circle your choice below.
ii) Explain the reasons for your choice. You must include reference to both acids in your answer.
iii) A sample of sodium ethanoate, CH3COONa, is dissolved in water. The solution is tested and found to be basic. Explain why the solution is basic.
Include appropriate equation(s) in your answer.
i) Explain why the concentration of the acid, HCl, in a 0.0376 mol L–1 HCl solution. is equal to the concentration of the hydronium ion, H3O+. 0.0376 mol L–1 HCl solution.
ii) The concentration of the hydronium ion, H3O+, in a 0.0376 mol L–1 solution of CH3COOH is less than 0.0376 mol L–1.
Explain why the concentration of the hydr onium ion is less than 0.0376 mol L–1.
iii) Conductivity of solutions can be described as being high, low, or having no conductivity. Compare and contrast the conductivity of the three solutions shown below.
0.100 mol L–1 HCl
0.100 mol L–1 CH3COOH
0.100 mol L–1 NaOH
Aqueous solutions of acids HA and HB both have the same concentration of 0.100 mol L–1. The pH of the solution of acid HA is 3.5 and the pH of the solution of acid HB is 1.8.
a) (i) Identify which one of these acids is stronger and circle your choice below.
ii) Discuss the reasons for your choice. You should include relevant equations in your answer, as well as reference to what is meant by the strength of an acid.
b) Describe what is observed when the following two tests are carried out on 5 mL samples of the acids HA and HB.
i) Identical small pieces of magnesium ribbon are placed in each acid.
ii) Sodium hydroxide solution is slowly added to each acid. The volume of sodium hydroxide solution required to completely react with the acid is measured.
c) Discuss the observations in (b) (i) and (ii).
Your answer must include reference to:
• similarities and / or differences in the observations of the tests on each acid
• equations for reactions.
Methyl orange can be used as an acid-base indicator. It is pink in solutions with a pH less than 3 and yellow in solutions with a pH greater than 4.
Four beakers are known to contain one each of:
• 0.1 mol L–1 HCl
• 0.01 mol L–1 HCl
• distilled water
• 0.1 mol L–1 NaOH
a) Write down the pH and colours of each of the four solutions if methyl orange is added to them.
b) Using only the methyl orange indicator, additional water, test tubes and a measuring cylinder, discuss how a student could identify each of the four solutions.
A solution of sodium hypochlorite, NaOCl, is basic. Discuss the above statement, including appropriate chemical equation(s) in your answer.
The concentration and pH of three acids, HA, HB and HC, are shown in the table below.
a) A small piece of magnesium is added to a 20 mL sample of each of the acids. Choose acid that would be expected to react most rapidly with the magnesium.
Explain why this acid will react the fastest.
b) Choose the weakest acid, Explain why this acid is the weakest
c) A solution of sodium ethanoate (NaCH3COO) is tested and found to have a pH of 8.50. Discuss why the pH of the solution is greater than 7. Include appropriate equation(s) in your answer.
Two acids of the same concentration, hydrochloric acid (HCl) and propanoic acid (CH3CH2COOH), have properties as shown below:
(b) Consider the properties described in the table above. Explain the differences in the conductivity and pH of the two acids. In your explanation include reference to the species present in each solution.
examiners tips: Read these please!
acid strength is the ability to donate protons (H+'s) to water
an increase in pH indicates a decrease in hydrogen concentration [H3O+] or [H+]
The table below provides an overview of the specify factor affecting the Rate of a Reaction and the year that this question appeared in the NCEA Exam paper
(a) Hydrochloric acid was reacted with calcium carbonate in the form of marble chips (lumps) and powder (crushed marble chips) in an experiment to investigate factors affecting the rate of a chemical reaction.
(i) Identify the factor being investigated.
(ii) Explain why the hydrochloric acid would react faster with the powder.
(b) A clock reaction involves mixing solution X and solution Y with starch present. When the reaction is complete the solution turns blue-black in colour.
A student carried out this reaction between solution X and solution Y in a conical flask. Over time, the cross on the piece of paper under the flask disappeared when viewed from above.
Elaborate on why the reactions in Experiment 2 and Experiment 3 occur faster than the reaction in Experiment 1. In your answer, include the following words or terms: collisions activation energy temperature effective catalyst
When dilute hydrochloric acid, HCl(aq), is added to sodium thiosulfate, Na2S2O3(aq), in a conical flask, the following reaction occurs:
2HCl(aq) + Na2S2O3(aq) → 2NaCl(aq) + SO2(g) + S(s) + H2O(ℓ)A pale yellow solid of sulfur, S(s), forms during the reaction. Over time, a cross on a piece of paper under the conical flask gradually disappears when viewed from above.
(a) List TWO ways that the rate of this reaction could be decreased.
(b) The following experiments were carried out, and the times taken for the cross to disappear recorded. The HCl(aq) was in excess in all of the experiments.
Analyse how the results of Experiment 2 and Experiment 3 compare to Experiment 1.
In your answer you should:
• identify the factor being changed and the effect it has on the reaction rate
• explain how the rate of reaction was affected, with reference to the collision of particles, and activation energy where appropriate.
Experiment 2 compared to Experiment 1:
Experiment 3 compared to Experiment 1:
The reaction between hydrochloric acid solution and sodium thiosulfate solution produces a precipitate of sulfur that makes the resulting solution go cloudy.
The reaction is carried out in a conical flask, which has been placed on a piece of paper with a cross, as shown in the diagram below.
The rate of this reaction can be measured by timing how long it takes for the solution to go so cloudy that the cross can no longer be seen through the solution.
The equation for this reaction is represented by:
2HCl(aq) + Na2S2O3(aq) --> S(s) + SO2(g) + 2NaCl(aq) + H2O(ℓ)(a) Describe how the rate of this reaction would change if the reaction was repeated:
(i) with a more diluted solution of sodium thiosulfate.
(ii) at a lower temperature.
(b) Explain how decreasing the concentration of sodium thiosulfate and decreasing the temperature of the reactants affects the reaction rate. In your answer, you should refer to:
particle collision
activation energy
the reaction between hydrochloric acid and sodium thiosulfate
(a) Zinc foil reacts with dilute sulfuric acid to produce zinc sulfate and hydrogen gas.
Zn(s) + H2SO4(aq) → ZnSO4(aq) + H2(g)Two experiments were carried out using the same mass of zinc foil and excess sulfuric acid of concentrations 0.500 mol L–1 and 1.00 mol L–1. The reactions were carried out at 25°C. A graph showing the results of the two experiments is shown below.
(i) Write down the letter of the experiment which was carried out using the 1.00 mol L–1 sulfuric acid. Justify your answer in terms of the collision of particles.
(ii) Experiment A was repeated, using the same mass of zinc granules instead of zinc foil. The rate of reaction increased. Explain this effect on the reaction rate, in terms of the factor investigated and the collision of particles.
(iii) Experiment A was repeated again; this time the reaction was carried out at a temperature of 15°C. Discuss the effect of this temperature change.
Your answer must include reference to:
reaction rate
collision of particles
Activation Energy
The experiment was repeated under different conditions as given below.
Complete the table below for EACH change, by:
stating the factor being investigated
describing how the reaction rate would be affected by using the word "increase" or "decrease" in your answer
(a) An experiment was carried out by reacting calcium carbonate with dilute hydrochloric acid.
The equation for this reaction can be represented by:
CaCO3(s) + 2HCl(aq) → CaCl2(aq) + CO2(g) + H2O(ℓ)State THREE ways to increase the rate of this reaction.
(b) Three experiments were carried out by reacting magnesium metal with hydrochloric acid.
The equation for this reaction can be represented by:
Mg(s) + 2HCl(aq) → MgCl2(aq) + H2(g)The beakers drawn below show particle representations of magnesium metal and hydrochloric acid just before the reaction is about to begin.
The rate of reaction can be followed by recording the volume of hydrogen, H2, produced. The graph below shows the volume of hydrogen produced during the initial part of the reaction.
Which beaker(s) corresponds to: LINE A of the graph and LINE B of the graph _
Explain your answers in terms of
particle collision
concentration of the particles present
(c) An experiment was carried out by reacting bromine, Br2, and methanoic acid, HCOOH, in a beaker. The equation for this reaction can be represented by:
Br2(aq) + HCOOH(aq) → 2Br–(aq) + 2H+(aq) + CO2(g)brown colourless all colourlessThe rate of reaction can be followed by measuring the time taken for the solution to go colourless.
A series of reactions were carried out and the time taken for the solution to go colourless was recorded. The results are shown below.
(i) Describe the trend shown by these results.
(ii) Discuss the reasons for this trend.
Your answer must include reference to
particle collision
reaction rate
Activation Energy
The rate of decomposition is changed with the addition of manganese dioxide, MnO2.
The rate of the reaction can be followed by recording the volume of oxygen produced. The graph below shows the volume of oxygen produced with, and without, manganese dioxide.
(a) Describe what happens to the rate of decomposition of hydrogen peroxide when manganese dioxide is added.
(b) A student suggested that manganese dioxide is a catalyst for the decomposition reaction.
Explain the role of a catalyst in changing the rate of a reaction.
Your answer must include reference to:
particle collision
Activation Energy
(c) One line on the graph above is divided into two time periods, A and B.
Discuss the change in reaction rate in terms of particle collisions during these two time periods, A and B.
An experiment was carried out by reacting zinc metal with excess dilute sulfuric acid. All the zinc was used up in the reaction. The equation for this reaction is:Zn(s) + H2SO4(aq) → ZnSO4(aq)+ H2(g)(a) The experiment was repeated, but this time 2 mL water was also added to the beaker. Again the zinc was all used up. Explain why: when water is added, the rate of reaction will decrease.
(b) Small pieces of copper can act as a catalyst for the reaction between zinc and dilute sulfuric acid.
Discuss how this claim could be tested by experiment.
In an experiment, a sample of large marble chips (CaCO3) is added to 200 mL of dilute hydrochloric acid in an open conical flask. The reaction that occurs is shown below.
2HCl(aq) + CaCO3(s) → CaCl2(aq) + H2O(ℓ) + CO2(g)As the carbon dioxide gas escapes from the flask, the total mass of the flask and contents decreases. The loss in mass is recorded at 5 minute intervals until the reaction has stopped. The experiment is repeated, using the same mass, but different sized, marble chips.
The results are shown in the table below.
(a) What factor affecting reaction rate is being investigated in this experiment?
(b) What conclusion can be made about the reaction rate using the data in the table?
(c) Explain the change in reaction rate that occurs, with reference to the collisions of particles.
(1) Hydrogen peroxide decomposes at room temperature (25'C) according to the following equation.2H2O2(aq) → 2H2O(l) + O2(g)
As the decomposition reaction occurs, bubbles of gas are produced.
(a) On addition of a very small amount of solid manganese dioxide, the rate at which the bubbles of gas are produced is increased so that rapid fizzing is observed. Further observation indicates that manganese dioxide remains after reaction has stopped. With reference to the collisions of particles, explain why the reaction rate has increased.
(b) Hydrogen peroxide is stored at a low temperature. Discuss this statement in terms of reaction rate.
(2) The concentration and pH of three acids, HA, HB and HC, are shown in the table below.(a) A small piece of magnesium is added to a 20 mL sample of each of the acid. Name the acid that would be expected to react most rapidly with the magnesium. Explain why this acid will react the fastest.
(a) How would the time taken for the solution to go cloudy be affected if the reaction were carried out in a water bath at a temperature of 50°C?
(b) With reference to the collisions of particles, explain why the reaction is affected in this way.
There is no set of Crystal Ball questions for Rates of Reaction, once you have completed all above you'll be very comfortable answering a question on rates of reaction.NCEA EXAM QUESTION:Write the Equilibrium constant expression for the following reactions
1) For the reaction below, the values of Kc at different temperatures are shown in the table. N2(g) + 3H2(g) ⇄ 2NH3(g)
Use this information to determine whether the formation of NH3(g) is endothermic or exothermic. Justify your reasoning using equilibrium principles.
2) For the reaction below, the Kc value is 46.8 at 491°C H2(g) + I2(g) ⇄2HI(g) Calculate the concentration of HI(g), at equilibrium, at 491°C, if the concentration of H2(g) is 0.0190 mol L–1 and the concentration of I2(g) is 0.210 mol L–1.
The table below shows the value of the equilibrium constant, Kc at two different temperatures.
(i) Circle the species that will be in the highest concentration at 200°C. PCl5(g) or PCl3(g)
(ii) Explain your answer.
(iii) Calculate the concentration of PCl5 at equilibrium at 350°C, if the concentrations of PCl3 and Cl2 are both 0.352 mol L–1.
The reaction between sulfur dioxide gas and oxygen gas at a particular temperature to make sulfur trioxide gas can be represented as:2SO2(g) + O2(g) --> 2SO3(g)For this equilibrium at this temperature, the value of Kc is 280.
The following three gases are mixed in a 1L container.
Justify whether the reaction would proceed in the forward OR reverse direction when the gases are mixed.
The following reaction can be used to produce gaseous methanol, CH3OH, from carbon monoxide and hydrogen.
CO(g) + 2H2(g) → CH3OH(g) ΔrH = –90.7 kJ mol–1At 25°C, the equilibrium constant, Kc = 2.20 × 10–4. Explain what this indicates about the relative amounts of reactants and product at equilibrium.
Hydrogen gas and iodine vapour are placed in a sealed container at 445°C. These gases combine to form hydrogen iodide gas. This equilibrium can be represented by:
H2(g) + I2(g) → 2HI(g)colourless purple colourless(a) Describe how an observer would know that the system had reached equilibrium.
(b) At 445°C, the equilibrium constant Kc = 49.5.
(i) Name the species that will be in the highest concentration at this temperature. You may name one, or more than one species in your answer.
(ii) Explain your answer.
(c) When the temperature of the equilibrium system is raised from 445°C to 1000°C (at constant pressure), the value of Kc decreases.
Use this information to determine whether the reaction between hydrogen and iodine is exothermic or endothermic. Justify your reasoning.
(a) Nitrogen monoxide gas reacts with oxygen gas to form nitrogen dioxide gas. The equilibrium reaction can be represented by:
2NO(g) + O2(g) → 2NO2(g)colourless colourless brownAt 230°C the equilibrium constant for this reaction has a value of 6.44 × 105.
(ii) State which gas will be in the highest concentration at 230°C. Explain your answer in terms of Kc and the colour seen.
Ag+(aq) + 2NH3(aq) → Ag(NH3)2+(aq)(i) At 25°C the value of Kc is 1.70 ´ 107. Name the species that would be present in the higher concentration in the equilibrium mixture at this temperature. Choose either Ag+(aq) or Ag(NH3)2+(aq) Justify your choice.2NO2(g) → 2NO(g) + O2(g)(ii) At 200°C the value of Kc is 1.10 ´ 10–5. Name the species that would be present in the higher concentration in the equilibrium mixture at this temperature. Choose either NO2(g) or NO(g). Justify your choice.NCEA EXAM QUESTION: Describing and Explaining equilibrium constant reactionsNC
NCEA EXAM QUESTION: Describing and Explaining equilibrium constant reactions
The two reactions shown in the following table are both at equilibrium.
Compare and contrast the effect of increasing the pressure on both reactions, with reference to the equilibrium positions.1) For each of the following changes applied to this system:
(i) State if the amount of chlorine gas, Cl2(g),would increase or decrease.
(ii) Justify your answers using equilibrium principles.
PCl3(g) is removed
The pressure is decreased.
2) When the temperature of the equilibrium system is increased from 200°C to 350°C (at constant pressure), the value of Kc increases, as shown in the table
Use this information to determine whether the decomposition of PCl5 is endothermic or exothermic. Justify your reasoning using equilibrium principles.
(a) Two oxides of nitrogen exist in an equilibrium system represented by the following equation. N2O4 is a colourless gas and NO2 is a brown gas.
N2O4(g) → 2NO2(g) ∆rH = + 57 kJ mol–1For each of the following changes applied to this equilibrium system in (i) and (ii) below:
describe the expected observation
use equilibrium principles to explain this observation.
(i) The reaction mixture is cooled.
(i) The pressure is decreased.
The following reaction can be used to produce gaseous methanol, CH3OH, from carbon monoxide and hydrogen.
CO(g) + 2H2(g) → CH3OH(g) ΔrH = –90.7 kJ mol–1Below are three possible factors that could be changed to alter the amount of methanol produced in this reaction.
Discuss the changes which could be made to increase the amount of methanol produced. Your answer should relate each of the factors in the table above to equilibrium principles.
The following equilibrium system is formed when hydrogen gas is mixed with iodine gas.
H2(g) + I2(g) → 2HI(g)colourless purple colourlessThe reaction has a negative value for ΔrH. For each of the following changes applied to this system:
describe the expected observation
use equilibrium principles to discuss the reason for this observation.
(i) HI(g) is added.
(ii) The reaction mixture is cooled.
(iii) The pressure is increased.
(1) One step in the production of sulfuric acid involves forming sulfur trioxide from sulfur dioxide. The equilibrium reaction can be represented by
2SO2(g) + O2(g) → 2SO3(g) ΔrH = –196 kJ mol–1(a) (i) Explain why a low temperature favours the formation of SO3(g).
(ii) The temperature that is actually used is approximately 450°C. However, this is not considered to be a
low temperature. Discuss why this temperature is used.
(b) (i) Describe another way of increasing the amount of SO3(g) present at equilibrium without adding any more reactants.
(ii) Explain why this will increase the amount of SO3(g) present at equilibrium.
(2) Carbon dioxide is added to drinks to make them fizzy. The following equilibria are involved:
The drink is fizzy when there is dissolved carbon dioxide, CO2(aq). The drink stops being fizzy when the carbon dioxide escapes from the drink as a gas.
Using equilibrium principles, discuss the changes that occur as a bottle containing fizzy drink is opened.
Your answer must include reference to:
equilibrium shift in Equation One and Equation Two
changes in the fizziness of the drink
any change in pH.
The following equilibrium system is formed when potassium thiocyanate solution is added to a solution of iron(III) nitrate. The reaction has a positive value for ΔrH
For each of the following changes applied to this system:
describe the expected observation
use equilibrium principles to discuss the reason for this observation.
(i) The reaction mixture is cooled.
(ii) Solid sodium fluoride is added to the reaction mixture. The fluoride ions react with Fe3+ ions.
(iii) Solid iron(III) chloride is added to the reaction mixture.
An equilibrium system is shown below.
3H2(g) + N2(g) --> 2NH3(g)The pressure of the system is increased, while maintaining a constant temperature. The percentage of NH3 in the reaction mixture is recorded and graphed.
(b) On the above graph, identify the line that shows the correct relationship between the percentage of NH3 in the reaction mixture, and increasing pressure. Explain your answer by applying knowledge of equilibrium principles.
(2) An equilibrium system involving different species of cobalt(II) is shown in the equation below.
At room temperature (25°C) the equilibrium mixture is pink.
(a) Describe the expected observation when solid sodium chloride (NaCl) is added to the equilibrium mixture. Explain your answer.
(b) The enthalpy change (∆rH) for this reaction as written above, has a negative value.
Name the ion that would be present in the higher concentration when the equilibrium mixture is heated, choose either [CoCl4]2–(aq) or [Co(H2O)6]2+(aq) Explain your answer.
The following reaction is exothermic:
Both N2O5 and O2 are colourless gases and NO2 is a brown gas. A mixture of these gases exists at equilibrium and is observed as a brown colour.
(a) For each of the following changes applied to the equilibrium system, describe the expected observation and explain why this occurs.
(i) The mixture of gases is heated (at constant pressure).
(ii) The pressure is increased, by decreasing the volume of the container.
(a) The following equilibrium system is established when thiocyanate ions (SCN) are added to iron (III) ions (Fe3+). The resulting aqueous solution is a dark red colour. The equation representing the equilibrium system and the colours of each species involved are given below.
Justify whether the reaction in which NH3 is formed, is endothermic or exothermic.
Complete the table showing the hydronium ion concentration, hydroxide ion concentration and pH for these solutions. Kw = [H3O+] [OH-] = 1.00 x 10-14
There is no Crystal Ball Activity for this task as once you have completed all the above calculations you will be sweet with them.(i) Complete the table below to show the conjugate acid-base pairs.
(ii) HPO42–(aq) is a species that can act as an acid or a base. Write equations for the reactions of HPO42– with water: one where it acts as an acid, and one where it acts as a base.
Complete the equations below to show how each species will react with water.
i) NH3 + H2O
ii) C2H5COO– + H2O
i) Identify one conjugate acid / base pair for the following reaction:
NH4+ + PO43– --> NH3 + HPO42–
ii) HSO4–(aq) is a species that can act as an acid or a base.
Write two equations for reactions of HSO4– with water: one equation where it acts as an acid, and one where it acts as a base.
Complete the equations below to show how each species will react with water to form an acidic solution.
i) NH4+(aq) + H2O -->
ii) HF(aq) + H2O -->
The bicarbonate ion, HCO3–, can both accept and donate hydrogen ions (protons). Write the equations below. When sodium bicarbonate, NaHCO3, dissolves in water the solution is basic.
Circle Reaction A or Reaction B to show which reaction predominates. Justify your answer.
i) Complete the table below to show the conjugate acid-base pairs.
ii) Circle the ion below that can act as both an acid and a base. Justify your choice CH3COO– HCO3–
iii) Hypochlorous acid is a weak acid. Complete the equation below to show the reaction of hypochlorous acid with water.
iii) A solution of sodium hypochlorite, NaOCl, is basic. Discuss the above statement, including appropriate chemical equation(s) in your answer.
Chickens make egg shell, CaCO3, using carbon dioxide gas from the air. The carbon dioxide forms carbonic acid (H2CO3), which then reacts to form the carbonate ions (CO32–) needed to make egg shell. Two equations showing part of this process are given below.
Equation 1: H2CO3(aq) + H2O(l) -->HCO3−(aq) + H3O+(aq)
Equation 2: HCO3−(aq) + H2O(l) --> CO32−(aq) + H3O+(aq)
(a) Identify three conjugate acid-base pairs in the equations above.
(b) HCO3− can act as both an acid and a base.
Specify which equation above (1 or 2) shows HCO3− acting as an acid. Give a reason for your answer.
NCEA PAST EXAM QUESTIONS on Explaining properties of aqueous solutions
Some properties of three aqueous solutions A, B and C, of equal concentration are shown in the table below.
The labels of the three solutions have been removed. It is known that the solutions are NH3(aq), HCl(aq) and NH4Cl(aq). Use the information in the table above to identify each of the three solutions. Justify the identification of all three solutions. In your answer you should:
• refer to both pH and electrical conductivity of the solutions
• link your answers to appropriate chemical equations.
(a) Place the following solutions in order of increasing pH,
0.01 mol L–1 CH3COOH
0.01 mol L–1 HCl
0.1 mol L–1 HCl
Justify your order above in terms of: proton transfer and relative concentrations of [H3O+] and [OH–], linked to the pH of the solution. You should include equations in your answer.
Three aqueous solutions, of equal concentration, have the following pH values
Compare and contrast both the strength and electrical conductivity of these aqueous solutions. Include appropriate equations in your answer.
i) The concentration and pH of two acids HA and HB are shown in the table below. Identify which one of these acids is weaker and circle your choice below.
ii) Explain the reasons for your choice. You must include reference to both acids in your answer.
iii) A sample of sodium ethanoate, CH3COONa, is dissolved in water. The solution is tested and found to be basic. Explain why the solution is basic.
Include appropriate equation(s) in your answer.
i) Explain why the concentration of the acid, HCl, in a 0.0376 mol L–1 HCl solution. is equal to the concentration of the hydronium ion, H3O+. 0.0376 mol L–1 HCl solution.
ii) The concentration of the hydronium ion, H3O+, in a 0.0376 mol L–1 solution of CH3COOH is less than 0.0376 mol L–1.
Explain why the concentration of the hydr onium ion is less than 0.0376 mol L–1.
iii) Conductivity of solutions can be described as being high, low, or having no conductivity. Compare and contrast the conductivity of the three solutions shown below.
0.100 mol L–1 HCl
0.100 mol L–1 CH3COOH
0.100 mol L–1 NaOH
Aqueous solutions of acids HA and HB both have the same concentration of 0.100 mol L–1. The pH of the solution of acid HA is 3.5 and the pH of the solution of acid HB is 1.8.
a) (i) Identify which one of these acids is stronger and circle your choice below.
ii) Discuss the reasons for your choice. You should include relevant equations in your answer, as well as reference to what is meant by the strength of an acid.
b) Describe what is observed when the following two tests are carried out on 5 mL samples of the acids HA and HB.
i) Identical small pieces of magnesium ribbon are placed in each acid.
ii) Sodium hydroxide solution is slowly added to each acid. The volume of sodium hydroxide solution required to completely react with the acid is measured.
c) Discuss the observations in (b) (i) and (ii).
Your answer must include reference to:
• similarities and / or differences in the observations of the tests on each acid
• equations for reactions.
Methyl orange can be used as an acid-base indicator. It is pink in solutions with a pH less than 3 and yellow in solutions with a pH greater than 4.
Four beakers are known to contain one each of:
• 0.1 mol L–1 HCl
• 0.01 mol L–1 HCl
• distilled water
• 0.1 mol L–1 NaOH
a) Write down the pH and colours of each of the four solutions if methyl orange is added to them.
b) Using only the methyl orange indicator, additional water, test tubes and a measuring cylinder, discuss how a student could identify each of the four solutions.
A solution of sodium hypochlorite, NaOCl, is basic. Discuss the above statement, including appropriate chemical equation(s) in your answer.
The concentration and pH of three acids, HA, HB and HC, are shown in the table below.
a) A small piece of magnesium is added to a 20 mL sample of each of the acids. Choose acid that would be expected to react most rapidly with the magnesium.
Explain why this acid will react the fastest.
b) Choose the weakest acid, Explain why this acid is the weakest
c) A solution of sodium ethanoate (NaCH3COO) is tested and found to have a pH of 8.50. Discuss why the pH of the solution is greater than 7. Include appropriate equation(s) in your answer.
Two acids of the same concentration, hydrochloric acid (HCl) and propanoic acid (CH3CH2COOH), have properties as shown below:
(b) Consider the properties described in the table above. Explain the differences in the conductivity and pH of the two acids. In your explanation include reference to the species present in each solution.
examiners tips: Read these please!
acid strength is the ability to donate protons (H+'s) to water
an increase in pH indicates a decrease in hydrogen concentration [H3O+] or [H+]
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